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Home , Borane, Catenation, Diagonal relationship, Hydride, Hypochlorite, Ionization energy

OF THE ELEMENTS 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

14-1 Chapters 14 and 23. PROPERTIES OF THE ELEMENTS First, let’s look at Periodic Table and numbering.

14.1 . The simplest , ~90% of all in the universe. The only element whose isotopes are each given a different symbol and name. 1H (or H, protium) = one (p+), plus one e- surrounding it. 2H (or D, ) = one p+ and one (n), plus one e- 3H (or T, ) = one p+ and two n, plus one e-.

Deuterium (2H) was produced in the ‘Big Bang’ — it is too fragile to survive fusion conditions in the stars (which produce the lighter elements) or supernovas (which produce heavier elements).

Hydrogen is the exception in the periodic table — it cannot be satisfactorily classified in any group: it has similarities both to (a) group 1 such as Li, Na, etc, in forming H+ and (b) group 17 non-metals such

as F, Cl, etc, in being H2 (H-H with a single ) in its stable - elemental form (compare F2, Cl2, etc) and also forming H ( ) analogous to F-, Cl-, etc.

14-2 Brief Summary of Hydrogen - most commonly forms covalent compounds; ionic compounds are rare. - high (unlike group 1, because e- close to nucleus without other e’s to shield it) and low (unlike group 1, because only one proton to attract e-s).  unlike groups 1 and 17 in that H+ and H- are rare (whereas Na+, K+, etc, ions are common, as are F-, Cl-, etc ions) because they usually bond + - + covalently to other things e.g., H3O , OH , NH4 , etc Very rare exceptions are certain ionic salts of H-, the hydride ion, in

compounds such as NaH and CaH2 (similar to NaCl and CaCl2).

Ionic (H-). With very strong reducing agents (Na(s), Ca(s), Li(s), etc.), hydrogen is reduced to H- = ionic hydrides.

e.g. 2 Na (s) + H2 (g)  2 NaH (s) - - Note that: H2 (g) + 2 e  2 H (g) Eº = -2.23 V (very negative!) Hydrides are thus very reactive (strong reducing agents) and will either: + (1) react with a H and go to H2 (g): + - NaH (s) + H2O (l)  Na (aq) + OH (aq) + H2 (g)

or (2) reduce something and go to H2 (g): TiCl4 (l) + 4 LiH (s)  Ti (s) + 4 LiCl (s) + 2 H2 (g) 14-3 Covalent Hydrogen Compounds. Very common and stable: CH4,NH3,H2O, HF, etc, etc. These other elements have higher electronegativity than H (H = 2.2, C = 2.5, N = 3.1, O = 3.5 F = 4.1)  we think of these as containing H+ and C4-,N3-,O2-,F-

oxidation states. e.g. F2 (g) + H2 (g) → 2HF (g)

H2 is a very important gas, for many reasons. For example:

N2 (g) + 3 H2 (g)  2NH3 (g) G < 0 (spontaneous) but very slow under normal conditions due to very strong NN  reaction run in industry at high T (~400 ºC) and pressure (250 atm) with an Fe catalyst to speed it up. This is called the Haber

process,andisthemainsourceofNH3 for the industry.

Metallic (Interstitial) Hydrides. H2 and H atoms can occupy space in-between the atoms of a . In particular, (Pd) has a

high affinity and can hold vast amounts (Pd:  935 times its volume = PdH0.5). Best thought of as a solution of the gas in the metal! Came to people’s attention during stories of late 1980’s.

Formation of Pd/H2 is used to purify H2 from gas mixtures.

14-4 Figure 14.2 A metallic (interstitial) hydride

Many transition metals form metallic (interstitial) hydrides,

in which H2 molecules and H atoms occupy the holes in the of the metal.

14-5 ACROSS THE PERIODIC TABLE Increasing: metallic character non-metallic character

ionic covalent oxides

basicity of oxides acidity of oxides

electropositivity electronegativity

ALSO: (i) Trends for a given element are affected by changes in its oxidation state (ii) For Periods 2 and 3, atom size decreases left-to-right Figure 14.1 Examples: Which is the most metallic element a) Li or Be b) Be or Mg c) Al or K

Which is the most acidic ? (i.e. dissolve in H2O to give acidic soln) a) Na2OorCl2O7 b) BaO or As2O5 Which is the stronger ? a) Al or Mg b) Na or K c) Br- or Cl- Which is the stronger ?

a) S or Se b) Br2 or Cl2 c) O2 or F2 Summary: The more metallic an element, the more basic is its oxide, the more ionic is its oxide, and the more electropositive (less electronegative) it is. 14-6 Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the 2 Elements.

Trends in , , and electronegativity across Period 2. 14-7 GROUP 1. ALKALI METALS (ns1) Found as M+ in (loss of ns1 ). Too reactive (strong reducing agents) to be found as the free metal. Prepared in industry by electrolysis of melted salts.

All well-studied except — radioactive, longest-lived isotope is 283Fr: half-life of only 21.8 min. Estimated only ~25 g on Earth at any one time

Properties. Soft, silvery metals at 20 ºC. All react vigorously with

to give H2 gas: the reactivity increases down group.

M (s) + H2O (l)  H2 (g) + MOH (aq) (video)

half-reactions: M (s)  M+ (aq) + e- - - H2O (l) + e  ½H2 (g) + OH (aq)

+ - overall: M (s) + H2O (l)  ½H2 (g) + M (aq) + OH (aq)

14-8 All react with O2 gas but products depend on metal:

2- Li gives Li2O (contains O i.e. normal oxide ion) 2- 2- Na gives Na2O2 (contains O2 =O-O i.e. peroxide ion) - - K, Rb, Cs give MO2 (contains O2 =O-O i.e. ion)

All M(s) must therefore be stored under inert oils to prevent reactions with air and water.

Note: Burning in air (rather than pure O2 gas) gives above products, but Li (and only Li) also reacts with N2 (g) to give the , Li3N, containing Li+ and N3-.

6 Li (s) + N2 (g)  2 Li3N

14-9 Important Question – WHY is Li different? The difference between Li and the rest + of group 1 (e.g. gives oxide with O2,reactswithN2, etc) is due to the small size of Li -leadstohigh charge-to-size ratio. This also causes Li compounds to have significant covalent character (i.e. LiCl is still fairly ionic but has noticeable covalent character, whereasNaClisveryionic).Li+ has too high a charge-to-size ratio (“charge ”) tobehappyascompletelyfreeLi+ ion, so it shares a bit with anions (i.e. some covalency).

*** For this and other reasons (see later), the top member of every group is significantly different from the rest ***

Note 1: The decreasing charge-to-size ratio down the group (i.e. Li+ >Na+ >K+ >Rb+ > + Cs ) explains the products on reaction with O2. The highest charge density (highest charge-to-size ratio) is Li+, and this forms the highest charge density O2- , 2- - whereas down the group the metals form the O2 (peroxide) and then the O2 (superoxide) salts, i.e. decreasing charge density ions. i.e., high charge density Li+ favors forming a salt with high charge density O2-,but down the group progressively lower charge density M+ favor formation of progressively lower charge density anions.

Note 2: Group 1 metals have very high 2nd Ionization Potential -means M2+ not possible to make (in stable form at 25 ºC). Of course, in places such as the surface of the sun, all sorts of other ions are possible.

14-10 GROUP 2. ALKALINE EARTH METALS (ns2) Occur in nature as M2+ ― too reactive to exist as free metals. (Ra) radioactive: 226Ra: half-life of 1599 years.

Again, Be slightly different from the rest due to high charge-to- size ratio (cf. Li in group 1): very hard metal and toxic. Others softer and not toxic. Also, Be compounds are mostly covalent, since Be2+ has too high a charge density to be happy as the free ion in ionic compounds – therefore it shares electrons with other atoms, decreasing its charge density (i.e. covalent bonding).

Be is unreactive to H2O; Mg reacts slowly with cold H2O, but fast with steam; others react vigorously with H2O. M (s) + 2 H2O (l)  M(OH)2 (aq) + H2 (g)

All react with O2 to give oxides MO except Ba, gives peroxide BaO2.

Mg reacts with N2 gas to ionic nitride, like Li. Others only at high T. 3Mg(s)+N2(g) → Mg3N2(s) (Mg + CO2) 1st and 2nd are both very easy  M2+ (no M+). 3rd ionization very difficult ― no M3+ 14-11 Diagonal Relationships

- similarities in properties between a period (row) 2 element and period 3 element, one group to the right.

Consider: Li vs Mg - similar size of Li+ vs Mg2+ (0.76 vs 0.72 Å) and similar properties (e.g., both give

with N2, salts with similar solubilities and thermal stabilities, and significant covalency in their bonds). The origin of these similarities is again the high charge-to-size ratios of Li+ and Mg2+.MgisbiggerthanLi,soMg2+ and Li+ end up being similar size.

14-12 Section 14.5. GROUP 13. (ns2np1) Metallic character decreases moving right, and we find that (B) is not a metal = “” or “semi-metal”. Al video ** Metalloid = element with props between metals and non-metals ** - B2O3 is acidic (cf., Li2O is basic, gives OH in water). Down group, metallic character increases - remainder of group 13 are metals but oxides of Al and Ga are amphoteric (see below), while those of In and Tl are basic. Compare:

B: B2O3 (s) + 6 NaOH (aq)  2Na3BO3 (aq) + 3 H2O(l) salt water

(cf. HCl (g) + NaOH (aq)  NaCl (aq) + H2O(l) )

in water, forms B(OH)3 (or H3BO3), known as boric acid. - + B(OH)3 + 2 H2O ⇌ B(OH)4 + H3O pKa = 9.25 Al, Ga: oxides react with bases as above, but also with . They are “amphoteric” (= can behave as both acids or bases)

Al2O3 (s) + 6 NaOH  2Na3AlO3 (aq) + 3 H2O(l) acid base salt water

Al2O3 (s) + 3 H2SO4 (l)  Al2(SO4)3 (s) + 3 H2O(l) base acid salt water In, Tl: oxides are basic and react only with acids.

In2O3 (s) + 3 H2SO4 (l)  In2(SO4)3 (s) + 3 H2O (l) Base acid salt water 14-13 PERIODIC TABLE OF THE ELEMENTS 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

14-14 Bonding: Also varies - ionic character increases down the group (because metallic character is increasing). B compounds are all covalent, Al are sometimes ionic and

sometimes covalent (e.g. AlCl3); the rest mainly ionic (but also affected by oxidation state).

AlCl3 is covalent!

Oxidation States Allform+3,butTlalsocanbe+1. ** In general, when two oxidation states are possible, lower one becomes more important down the group, and its properties are more metal-like **

Again, B (period 2) more different from the rest of group. For example, B forms many electron-deficient compounds – these are stable but nevertheless react with Lewis bases ⇨ attain an octet.

- e.g., BF3 :F: Bhasonly6e in its outer () shell :F:B:F:  strong Lewis acid  will accept from Lewis base

e.g., F3B+:NH3  F3B─NH3

- + this is source of acidity of B(OH)3 +H2O  B(OH)4 +H

14-15 Uses: Plentiful. B2O3 used in production of borosilicate . B(OH)3 (boric acid) used as , eyewash, insecticide. Na2[B4O5(OH)4]8H2Ousedinwashing powders.

Boron hydrides BxHy very important class of compounds. Figure 14.8

Aluminum used in water purification, dye industry, antiperspirants, etc.

Al2O3 used as a support for industrial catalysts, chromatography supports, etc.

BN compounds similar to analogous C compounds. e.g. borazene, like benzene.

Borazon (BN) has a similar structure to diamond and thus also very hard.

Diagonal Relationships: Be with Al (and B with Si)

Be and Al have similar properties e.g., similar covalency in their compounds.

14-16 Figure 14.8 Some boron hydride structures (“”)

Borane is BH3

Diborane

14-17 Section 14.6. GROUP 14. (ns2np2) - first group to show complete range of properties, from non-metal (C)tometal(SnandPb)–showsupinpropertiessuchasmelting

points and Hfus - non-metals give strong covalent bonding  higher melting points (Table 14.2)

Elements: show important allotropes for the first time in C

chemistry: graphite (2D sheets), diamond (3-D network), Cn (molecular fullerenes, e.g., C60,C70, …), (single sheet of graphite). C60 looks like a soccer ball. Figure 14.10 e.g., diamond graphite

very hard soft and greasy (used as lubricant) colorless black 3D 2D sheets insulator conductor

Graphite is the standard state of C at 298 K and 1 atm. Diamond formed at high T and P and interconverts to graphite at 298 K (1 atm, but very slowly).

14-18 Allotropes of

Diamond Graphite graphene – single sheet of graphite

multi-walled nanotube C60 C70 single-walled nanotube

14-19 Oxidation states: ** multiple ones now more common! **

C: non-metal: compounds all covalent except C4- (carbide) ion e.g.

Ca2C. Oxide (CO2) acidic: almost all oxid. states from -4 to +4 (e.g., CH4 to CO2)known. Si, Ge: – essentially all compounds are +4 ox. state, but a few are +2.

Sn, Pb: metals +4 (covalent) e.g., MO2 +2 (ionic usually) e.g., MO Figure 14.11

C usually four-coordinate (exception CO). Other elements show more 2- 2- 2- 2- exceptions (e.g., SiF6 ,GeCl6 ,Sn(OH)6 ,Pb(OH)6 ) but still usually four-coordinate.

Note difference: Chas2s22p2 outer configuration with no available d orbitals. Remainder of elements have d orbitals they can use in bonding  can form six bonds (sp3d2 hybrids).

14-20 Figure 14.11 Salt-like M2+ and oily M4+ chlorides show greater metallic character of and in the lower oxidation state.

M2+ – , typical of metals M4+ – covalent bonding, typical of non-metals

14-21 Again: big difference between C and rest of group. Main ones: (i) ability to form lowest (-4) and highest (+4) oxidation states

(ii) forms multiple bonds to itself or other light elements (N, O) e.g. CO2 (i.e. O=C=O), acetone (i.e. (CH3)2C=O) (iii) forms stable/common single bonds to itself (“catenation” = chain formation).

e.g. n-butane is H3C-CH2-CH2–CH3 containing a C-C-C-C chain. Include O, N, S, etc. and you have organic and , and Life!

-bonding: C is small, forms strong enough  bonds to give stable compounds.

Elements further down the group: -bonding is much weaker and usually prefers to form extra sigma (σ) bonds.

Compare CO2 vs SiO2 both are oxidation state = +4 but 3D

strong C-O  bond weak Si-O  bond, therefore each  C=O is 1 sigma, 1 pi Si forms four Si-O bonds  each C has 2 sigma, 2 pi

14-22 C can even form further  bonding i.e., C≡O (1 sigma, 2 pi bonds)

CO2 (g) + C (s) ⇌ 2CO(g) H>0,S>0

Non-spontaneous (G > 0) at R.T but spontaneous (G<0)athighT.

In fact, favored at >700 ºC unless excess O2 present  CO is a major pollutant from processes at high T and/or in O2- starved conditions when 2 C + O2  2CO

Note: CO2 is acidic, CO is not

 CO2 (g) + H2O(l) ⇌ H2CO3 (aq) carbonic acid

but CO (g) + H2O(l) noreaction!

CO (g) very poisonous – binds to hemoglobin irreversibly.

14-23 (SiO4)n tetrahedra form many structures by condensation:

1D include .

3D polymers (SiO2), include , mica, feldspar, sand, clay, , gems, etc. Fig 14.15

Diagonal Relationship: B with Si

14-24 Figure 14.15 Quartz is a three-dimensional framework silicate.

14-25 Section 14.7. GROUP 15. (ns2np3) N, P = non-metals; As, Sb = metalloids; Bi = metal

N again different from the rest of the group.

All show +5 ox. state; +3 becomes more important down group

N2 gas (N≡N) vs. P4 or

1,2 bonds (3σ) (white P) (red P)

Oxidation States N all from +5 to -3 common P, As, Sb +5 and +3 common, others rare Bi only +3 is common, others very rare

Lowest ox. state (-3) reacts with H2O to give EH3 - toxic except NH3. Ca3As2 (s) + 6 H2O(l) 2AsH3 (g) + 3 Ca(OH)2 (aq)

N, P oxides acidic; As, Sb, amphoteric; Bi basic

Acidity of oxides increases with higher oxidation states

14-26 The oxides of N – many are known and important

N2ONON2O3 NO2 N2O4 N2O5 +1 +2 +3 +4 +4 +5 nitrous nitric dinitrogen dinitrogen dinitrogen oxide oxide trioxide dioxide tetraoxide pentoxide

(N2O is an anesthetic and ‘laughing gas’)

In contrast, only two N-containing oxo-acids are known: (must learn the –ous vs –ic and –ite vs –ate naming rules!!)

- HNO2 is (NO2 = ion) (N is +3 oxid state)

- HNO3 is (NO3 = ion) (N is +5 oxid state)

-4 HNO3 is a strong acid, HNO2 is weak (Ka =7.1x10 ,pKa=3.15)

14-27 Similarly for P (the element is spelled -not phosphorous!)

P4 (s) + 3 O2 (g)  P4O6 (s) oxid. st. +3

P4 (s) + 5 O2 (g)  P4O10 (s) oxid. st. +5

1) P4O6 +H2O  H3PO3 (phosphorous acid)

- 2- 3- diprotic acid  [H2PO3]  HPO3 (no PO3 ) (H bound to P cannot dissociate)

2) P4O10 +H2O  H3PO4 ()

- 2- 3- triprotic acid  H2PO4  HPO4  PO4

14-28 Section 14.8 GROUP 16. (ns2np4) O, S, Se non-metals Te metalloid Po metal (radioactive)

O is too electronegative to show maximum oxidation state + 6. S does, then lower oxidation states again become more important. Main oxidation state: O(-2), S(+6), Se(+4), Te(-2).

O found everywhere, e.g., as O2 and H2O. S also found as element, S8, others as minerals.

O2 is O=O (1 σ,1) gas. S occurs as yellow S8 solid (only σ)andother Sn. Similarly for Sen.

Oxides SO2 and SO3 give H2SO3 and H2SO4 in water (+4) (+6) (sulfurous acid)()

SeO2 and SeO3 give H2SeO3 and H2SeO4 (+4) (+6) ()(selenic acid)

14-29 Hydrogen compounds

H2O, H2S, H2Se, H2Te are all -2 oxidation state

Note: usually b.pts of compounds increase down a group but due to strong O-H⋯O hydrogen-bonding, b.pt. of liquid H2Oisveryhigh. S-H⋯S etc. are much weaker.

(Similarly for Group 15 hydrogen compounds (NH3,PH3,etc,but there they are all gases).

The high b.pt. of water (100ºCat1atm)isthereasonthatlifeon this is the way it is. Liquid water allowed a solvent for life’s chemistry to develop in.

Note:H2S, H2Se, H2Te very toxic. H2O obviously is not.

14-30 Figure 14.21 The dehydration of sugar by sulfuric acid.

14-31 Section 14.9. GROUP 17. The (ns2np5)

F = the most electronegative element  only -1 oxidation state. F2 is incredibly reactive – powerful oxidizing agent. Others show range from maximum (+7) to minimum (-1), with 1themostimportant.

All are non-metals  acidic oxides, dissolve in water to give acids

• Oxides:(+1)Cl2O (rare +4) ClO2 (+7) Cl2O7 Fig 14.25 acid anion ox. st.

HClO4 +7 HClO3 chloric acid +5 HClO2 +3 HClO +1 HCl -1

HBrO4 perbromic acid +7 HBrO3 bromic acid +5 HBrO2 bromite +3 HBrO +1 HBr hydrobromic acid bromide -1 Similarly for I 14-32 Figure 14.25 oxides.

lone e-

Cl2O

dichlorine heptaoxide

Cl2O7

ClO2 14-33 Figure 14.26 Crystals of tetrafluoride (XeF4).

14-34